Electrochemistry/Redox

Redox

Oxidation and reduction reaction. LEO the lion goes GER.

Oxidation = losing electrons

Reduction = gaining electrons

Oxidizing agent

helps another species oxidize/lose electrons, usually it itself gets reduced

Reducing Agent

helps another reduce/gain electrons, usually it itself gets oxidized

Oxidation State vs. Oxidation Number

Its the charges assuming ionic. But the state is writing out charges with sign first, then number after.

Number is with the roman numeral.

Rules for writing oxidation number?

Write the higher rules first in importance- so break lower rules before higher rules

Oxidation State Rule 1

Elemental state of atoms = charge of 0. eg. diatmic molecules.

Oxidation State Rule 2

Simple ions are the charge they’re written as. eg. Na⁺ has a charge of +1, and Cl^- has a charge of -1.

Oxidation State Rule 3

Sum of oxidation states = overall charge. The individual components must add up to the overall charge.

Oxidation State Rule 4

Alkali metals always have a charge of +1, Alkaline earth metals always have a charge of +2

Oxidation State Rule 5

Hydrogen has a charge of +1, but there can be exceptions, like in metal hydrides, where it can be -1.

Oxidation State Rule 6

Oxygen has a charge of -2. But sometimes it can be different, like -1 depending on molecule.

Balancing Redox Equations: Half Equations Method

Split into half-equations, balance each half cell with MAJOR OH, balancing major elements before using H₂O and H⁺ to balance out the hydrogens and oxygens. Use electrons to balance charges. Then combine back together using LCM of electrons.

Disproportionation reactions

When one species undergoes both oxidation and reduction. (check notes)

Just write out the species that undergoes both into 2 half equations, and then balance using MAJOR OH, then e^-, then LCM of e^-

What to do if the redox equation says acidic or basic?

If the equation basic, add OH to neutralize the H⁺ added earlier to balance it. Remember that OH and H make water, so water is formed.

How to Balance Using Oxidation Number?

First balance elements. Write out oxidation states. Match the changes in ON using LCM. The LCM is multiplied for each section for each part that has the change in ON. Then add H₂O and H⁺to balance O and H. Check notes.

Note about ON redox method

The oxidation state is per atom, so if the oxidation takes into account all number of molecules there are. Check notes.

Table 19 Organization

Bottom of page, left side= best at reduction, strong oxidizing agent.

Top of page, right side = best at oxidizer, strong reducing agent.

How to find overall voltage/ E° cell.

1) State species present

2) Start from the bottom left, find the species best at reducing. Then start at top right to find species best at oxidizing. Make sure you write it so the strong reducer and oxidizer are both on right hand side.

3) Balance using LCM of electrons.

4) Add together E°. When sign is reversed, the value of E° changes to positive if neg, and neg if pos. This is the overall voltage

Why do we add to find overall voltage?

Voltage is a ratio, so even if its scaled from multiplying by the LCM, the ratio stays the same, so addition to find overall.

At what E° is it spontaneous?

When E° is positive the reaction is spontaneous, when negative, non-spontaneous.

Spontaneous is when a strong oxidizer and reducer react, weak reducer and weak oxidizer, non-spontaneous.

When both are oxidizers/both reducers, no reaction.

What does it mean if ON increases or decreases?

ON increase = oxidation

ON decrease = reduction

What happens if the oxidation state is already 0, but there are other elements?

They can have states of 0 if everything else already adds up to its overall charge

Just from looking at table 19, how can you tell if the reaction will spontaneous?

will be spontaneous when oxidizing agent is lower than the reducing agent Won’t be spontaneous when the oxidizing agent is higher than the reducing agent

Voltaic Cell/Battery what are they?

Spontaneous reaction produces electricity using voltage and current

Standard hydrogen electrode and drawing

Standard hydrogen electrode (SHE), use an inert electrode, like Pt and C, and pump in hydrogen gas onto electrode, in a solution of H+ ions. Make sure to write it’s at standard conditions. 1 bar, 25 298 kelvin, 1 M

Draw a battery,

could be a with a SHE connected to another electrode surrounded in a solution of its ions

How are voltages considered positive or negative?

Hydrogen, or H⁺ is chosen as a standard, and others are positive or negative based on how well they can reduce in comparison to H⁺

So if the element is worse at reducing, then its negative, if the element is better at reducing than its positive

AN OX CARED

anode is the site of oxidation, cathode is the site of reduction. Anode mass decreases, cathode mass increases. Flow of electrons is from anode to cathode.

Cathodes only need to be the site of reaction, a surface to occur on, so it doesn’t need to be something specific

What is used to maintain neutrality?

Salt bridges connect the 2 solutions to each other and maintain electroneutrality for both half cells in redox. Cations move to cathode thru salt bridge, anions go to anode thru salt bridge

How to calculate E° not for voltaic cells

same as calculating voltage- find best reducer, find best oxidizer, flip equation, and add the E° voltages together to get E° cell

What to put on a drawing of voltaic cells?

Electrode surrounded by solution, salt bridge connecting its solutions together, flow from electrons from anode to cathode. Direction in which ions are produced in solutions. If oxidation, arrow from electrode to solution. If reduction from solution to electrode

What is rusting?

iron or Fe, reacting with H₂O and O₂ . It basically goes from its oxidized state to reduced state.

Ways to prevent rust?

1) Add paint layer to physically separate Fe from O₂/H₂O

2) Cathodic protection

What is cathodic protection?

When you surround Fe with a cheaper metal that oxidizes before Fe, like Aluminum, so Fe stays reduced as cathode. Stops rust

Redox Titration

titrations to find some unknown concentration, and redox and electron transfers

MonO4 is an oxidizing agent

MnO4 is reducing, so it oxidizes Fe+3 to Fe+2 even though the Fe+2 is a worse oxidizer than Fe, it aqueous and better for titrations. Memorize MnO4 equation

Steps for a MnO4 redox titration

Iron determination- do the half cell reactions to find the balanced redox equation with MnO4 and iron

What to do with the balanced equation after the titration?

Use the amounts from the titration, and the stoich ratio from the balanced redox equation.

Note about MnO4

it’s self indicating, and the color is rlly thick and strong purple color, so read from the top of the meniscus. Once endpoint reached, when reacting with Fe+2, becomes colorless

Using I- as a reducing agent

I- oxidizes, and it can’t titrate anything above it. Could be titrated with bleach, NaOCl, as it reacts with the OCl-, forming I2. You can see it forms I2 from the balanced redox equation

2I^- → I₂ + 2e^-

OCl^-+ H⁺ + 2e^- → Cl^- + H₂O

Step 2 of using I- as a reducing agent?

Titrate the I2 produced with thiosulfate, or S2O3^-2 to form I-. Starch is used as indicator. Find the balanced redox reaction of thiosulfate and I-. Then use the stoich ratio from the first and second balanced equations used. Write out equations.

How to determine endpoint of reactions?

Add starch indicator, since starch, I2 forms a deep blue color

Why do you need to make sure KI and H+ are in excess?

In the first reaction with OCl-, you need to make sure all the OCl- is reacted, so add in excess KI and H+

What happens if a rxn shifts left in a voltaic cell?

Note: Just look to see if shifting gives more of less of the desired product

If the sign of the rxn is negative, shifting left reduces the potential, increasing E cell. If the sign is positive, decreases the potential, decreasing E cell

What happens if a rxn shifts right in a voltaic cell?

Note: Just look to see if shifting gives more of less of the desired product

If the sign of the rxn is negative, shifting right decreases potential, decreases E cell. if sign of rxn is positive, shifting right increases potential, increasing E cell

Note about drawing voltaic cells for gases?

Use an inert electrode, like Pt, and funnel gas into electrode

What does voltmeter read at equilibrium?

At equilibrium, voltmeter would read 0v

Does changing the size of the electrode impact voltage?

No.

Why would the voltage potential in a voltaic cell become 0.00v?

The concentrations of what is being oxidized and reduced reach a point where oxidation and reduction potentials become equal, but opposite in sign

Memorize from redox titrations:

memorize the balanced equations for I^- S₂O_{3^{-2 ,}} MnO₄ ^-

Iodide vs Iodine

Iodine = I2, Iodide = I-